Chemistry Fundamentals: Units, Measurement, Atomic Theory, and the Mole

Vocabulary flashcards covering measurement systems, SI units, precision, accuracy, significant figures, states and classification of matter, atomic laws, subatomic particles, isotopes, ions, atomic mass, and mole concepts.

Meter (m)

($1\,\text{yard} = 36\,\text{inches}$ , while $1\,\text{meter} = 39.37\,\text{inches}$).

Kilogram (kg)

$1\,\text{kg}$ is equal to $2\,\text{lb}\,3\,\text{oz}$, and $1\,\text{gram}$ ($\text{g}$ ) is $\frac{1}{1000}\,\text{kg}$.

Mass vs. Weight

Mass is a measure of the quantity of matter within an object, whereas weight is a measure of the gravitational pull on that matter.

Volume

The amount of space occupied by a substance; its SI unit is $\text{m}^3$, but common lab units include liters ($\text{L}$), milliliters ($\text{mL}$), and cubic centimeters ($\text{cm}^3$), where $1\,\text{mL} = 1\,\text{cm}^3$.

Accuracy vs. Precision

Accuracy refers to how close a measured value is to the actual or true value, whereas precision refers to how close a series of measurements are to one another or how reproducible they are.

Exact Numbers

Numbers that have an unlimited number of significant figures because they come from exact counting of discrete objects, integral numbers in equations, or defined conversion factors.

Conversion Factor

A fractional ratio derived from a unit equation (e.g., $2.54\,\text{cm} = 1\,\text{in}$) with the given unit on the bottom and the desired unit on top, used to convert measurements.

Classification of Matter

The categorization of matter by composition into pure substances (elements or compounds) or mixtures (heterogeneous or homogeneous).

Pure Substance

A type of matter made up of only one component whose composition is invariant.

Mixture

A substance composed of two or more components in proportions that can vary from one sample to another.

Element

A pure substance composed of a single type of atom that cannot be chemically broken down into simpler substances.

Compound

A pure substance composed of two or more different elements in fixed, definite proportions.

Heterogeneous Mixture

A mixture in which the composition varies from one region to another, making different portions identifiable (e.g., wet sand).

Homogeneous Mixture

A mixture with a uniform composition throughout because its components mix uniformly at the atomic or molecular scale (e.g., tea with sugar).

Law of Conservation of Mass

A fundamental chemistry law stating that matter is neither created nor destroyed in a chemical reaction.

Law of Definite Proportions

A law stating that all samples of a given compound have the same proportions of their constituent elements by mass.

Law of Multiple Proportions

A law stating that when two elements (A and B) form two different compounds, the masses of element B that combine with a fixed mass of element A can be expressed as a ratio of small whole numbers.

Dalton's Atomic Theory

The theory stating that elements are composed of tiny, indestructible atoms; all atoms of a given element share identical mass and properties; atoms combine in simple whole-number ratios to form compounds; and atoms rearrange without changing identity during chemical reactions.

Atomic Number (Z)

The number of protons in an atom's nucleus, which identifies the specific chemical element.

Mass Number (A)

The total number of protons ($p$) plus neutrons ($n$) in an atom's nucleus ($A = p + n$).

Isotopic Notation (${}_Z^A \text{X}$)

A chemical symbol notation where $X$ is the chemical symbol, $A$ is the mass number in the upper left superscript, and $Z$ is the atomic number in the lower left subscript.

Ion

A charged particle formed when a neutral atom or molecule gains or loses electrons during a chemical change.

Cation

A positively charged ion formed when an atom loses one or more electrons (e.g., $\text{Na}^+$).

Anion

A negatively charged ion formed when an atom gains one or more electrons (e.g., $\text{F}^-$).

Atomic Mass

The average mass of an element's naturally occurring isotopes, weighted according to the natural abundance fraction of each isotope.

Mole (mol)

The SI unit for amount of substance, defined as the amount of material containing $6.02214 \times 10^{23}$ particles (Avogadro's number), equal to the exact number of atoms in $12\,\text{g}$ of pure carbon-12.

Avogadro's Number

The number of particles in one mole of any substance, equal to $6.02214 \times 10^{23}$.

Avogardos number conversion

1Mol=6.022×10²²

F° to C

C° = F° -32 / 1.8

C° to K°

C° + 273.15

Molar Mass

The mass in grams of one mole of an element's atoms or compound's molecules, which is numerically equal to its atomic or formula mass in atomic mass units ($\text{amu}$).

Mass-to-Atom Conceptual Plan

The two-step conversion process converting grams of an element to moles using molar mass, and then converting moles to number of atoms using Avogadro's number.

Plan for finding mols and atoms

G element (g/mol) → mol element (avocados #) → number of atoms

If atoms is in mL how do we convert?

Use density - g/mL

Molar mass

G/Mol